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Complexation of Ni(II) and Co(II) with 1,4-Dioxa-7,10,13-triazacyclopentadecane (L). Crystal Structure of [NiLCl][NiL(H2O)](ClO4)3 and Macrocycle-Induced Dioxygen Binding Carla Bazzicalupi, Andrea Bencini, Emanuela Berni, Antonio Bianchi,* Claudia Giorgi, Piero Paoletti,* and Barbara Valtancoli Department of Chemistry, University of Florence, Via Maragliano 75/77, 50144 Florence, Italy
Ni(II) and Co(II) coordination with the macrocyclic ligand 1,4-dioxa-7,10,13-triazacyclopentadecane (L) has been investigated in solution (Co(II) and Ni(II)) and in the solid state (Ni(II)). Complexation equilibria have been studied by means of potentiometric titrations in 0.15 mol dm-3 NaCl, at 298.1 ( 0.1 K, evidencing the formation of [ML]2+, [MLOH]+, and [ML(OH)2] (M ) Co, Ni) species. Under aerobic conditions, the mononuclear Co(II) complex binds dioxygen, forming the dibridged [Co2L2(O2)OH]3+. This is the unique oxygenated species formed. The importance of the bridging OH- anion for the formation of the oxygenated complex has been evidenced by competitive coordination of SCN-. Structural information has been obtained by resolving the crystal structure of the [NiLCl][NiL(H2O)](ClO4)3 compound. The structure contains [NiLCl]+ and [NiL(H2O)]2+ cations in which the metal ions have octahedral coordination environments determined by five ligand donor atoms and one Cl-/H2O. Introduction Synthetic macrocyclic ligands are efficient receptors for metal ion binding forming metal complexes characterized by high thermodynamic and kinetic stability.1 Similar complexes are particularly suitable for the study of successive interaction with further coordinating species, in particular when such species have reactive properties such as molecular dioxygen.2 It has been known for a long time that synthetic metal complexes are able to bind reversibly dioxygen in a manner related to the natural systems.3 This observation stimulated the synthesis of specific ligands1 and the development of a considerable amount of investigations.2 In this respect, cobalt, iron, copper, and manganese metal complexes are the most studied systems. In addition to their similarities with natural systems, such complexes exhibit interesting ability in the activation of molecular oxygen toward a number of oxidation reactions.2,4 Three main categories of synthetic ligands1 have been employed to obtain metal complexes displaying dioxygen binding properties and dioxygen activation: (i) cofacial porphyrin dimers,5 (ii) ligands with “protected” coordination sites, such as “lacunar”, “picket fence”, and “capped” molecules,6 and (iii) simple macrocyclic ligands.2 Regarding the third class of ligands, polyamino- and polyaminopolyoxamacrocycles have been the most studied compounds able to form metal complexes with interesting activity toward dioxygen.2 In particular, several studies have been devoted to dioxygen binding by Co(II) complexes with mononucleating macrocycles.7-14 In the present paper we report the results of a thermodynamic study regarding the uptake of molecular dioxygen to the Co(II) complex of the macrocyclic ligand * To whom correspondence should be addressed. Phone: +39-055-354845. Fax: +39-055-354845. E-mail: bianchi@ chim1.unifi.it.
1,4-dioxa-7,10,13-triazacyclopentadecane (L) in an aqueous solution, obtained by performing potentiometric measurement under anaerobic and aerobic conditions. A structural characterization of the ligand coordination properties in the solid state has been obtained by resolving the crystal structure of the [NiLCl][NiL(H2O)](ClO4)3 complex.
Experimental Section Synthesis of the Compounds. 1,4-Dioxa-7,10,13triazacyclopentadecane (L) was prepared as previously reported15 and used as L‚3HClO4. Crystals of [NiLCl][NiL(H2O)](ClO4)3 suitable for X-ray analysis were obtained by slow evaporation at room temperature of an aqueous solution containing L and NiCl2, in equimolar concentrations, at pH 10, upon addition of excess NaClO4. The compound gave satisfactory elemental analysis. Perchlorate salts of metal complexes with organic ligands are potentially explosive; these compounds must be handled with caution! X-ray Structure Analysis. A green prismatic crystal of [NiLCl][NiL(H2O)](ClO4)3 was mounted on an Enraf Nonius CAD4 X-ray diffractometer, which uses an equatorial geometry. Graphite-monochromated Mo KR radiation was used for cell parameter determination and data collection. A summary of the crystallographic data is reported in Table 1. Cell parameters were determined by least-squares refinement of diffractometer setting angles of 25 carefully centered reflections. The crystal of the compound belongs to the monoclinic family, space group P21/n (Z
10.1021/ie000095d CCC: $19.00 © 2000 American Chemical Society Published on Web 10/02/2000
Ind. Eng. Chem. Res., Vol. 39, No. 10, 2000 3485 Table 1. Crystal Data and Structure Refinement for [NiLCl][NiL(H2O)](ClO4)3 empirical formula formula weight temperature wavelength crystal system, space group unit cell dimensions
volume Z calculated density absorption coefficient crystal size final R indices [I > 2σ(I)]a R indices (all data)a a
C20H48Cl4N6Ni2O17 903.86 298 K 1.54178 Å monoclinic, P21/n a ) 9.071(8) Å b ) 12.217(3) Å c ) 31.990(5) Å 3531(3) Å3 4 1.700 mg/m3
R ) 90° β ) 95.18(3)° γ ) 90°
4.817 mm-1 0.3 × 0.25 × 0.1 mm R1 ) 0.0584, wR2 ) 0.1379 R1 ) 0.1370, wR2 ) 0.1660
R1 ) ∑||Fo| - |Fc||/∑|Fo|; wR2 ) [∑w(Fo2 - Fc2)2/∑wFo4]1/2.
) 4) with a ) 9.071(8) Å, b ) 12.217(3) Å, c ) 31.990(5) Å, β ) 95.18(3)°, and V ) 3531(3) Å3. Intensities of two standard reflections were monitored during data collection to check the stability of the diffractometer and of the crystal: no loss of intensity was recognized. A total of 3879 reflections, up to 2θ ) 100°, were collected. Intensity data were corrected for Lorentz and polarization effects, and an absorption correction was applied once the structure was solved by the Walker and Stuart method16 (max/min correction for µ and φ ) 1.094181/0.917339 and max/min correction for θ ) 1.039889/0.807881). The structure was solved by the direct method of the SIR9217 program. Refinement was performed by means of the full-matrix least-squares method of the SHELXL9318 program which uses the analytical approximation for the atomic scattering factors and anomalous dispersion correction from ref 19. The function minimized was ∑w(Fo2 - Fc2)2, with w ) 1/[σ2(Fo2) + (aP)2 + bP], where a and b are refined parameters and P ) (Fo2 + 2Fc2)/3. Anisotropic displacement parameters were used for all of the non-hydrogen atoms. The hydrogen atoms were introduced in the calculated position, and their coordinates were refined in agreement with those of the linked atoms, with overall refined thermal parameters for the aliphatic and aminic hydrogens, respectively. A double position was found for the C11 atom and introduced in the calculation (C11 and C11′) with population parameter 0.5. For 457 refined parameters, the final agreement factors were R1 ) 0.0584 (for 2156 reflections with [I > 2σ(I)]) and wR2 ) 0.1660 (all data). Potentiometric Measurements. All pH-metric measurements (pH ) -log [H+]) were carried out in 0.15 mol dm-3 NaCl, CO2-free solutions at 298.1 ( 0.1 K, by using the equipment and the methodology that have been already described.20 The combined Ingold 405 S7/ 120 electrode was calibrated as a hydrogen concentration probe by titrating known amounts of HCl with CO2free NaOH solutions and determining the equivalent point by Gran’s method,21 which allows one to determine the standard potential E° and the ionic product of water (pKw ) 13.73(1) at 298.1 ( 0.1 K in 0.1 mol dm-3 NaCl).
Figure 1. ORTEP25 view of the [NiLCl]+ and [NiL(H2O)]2+ cations in [NiLCl][NiL(H2O)](ClO4)3.
Measurements performed to determine the equilibrium constants for the formation of the dioxygen complexes were carried out by maintaining a constant pressure (1 atm) of O2 into the potentiometric cell. The concentration of O2 in solution (1.21 × 10-3 mol dm-3), at the operating pressure, temperature, and ionic strength, was obtained from published data.22 A total of 15 min was allowed to elapse after each titrant (NaOH solution) addition to ensure the attainment of equilibrium conditions for all measurements. At least three potentiometric titrations were performed for each system in the pH range 2.5-11. Ligand concentration [L] ) 1 × 10-3 mol dm-3 and metal ion concentration [M] ) 0.8[L] were adopted in the complexation experiments under both anaerobic and aerobic conditions. The coordination of SCN- in the system M/L/SCN- was studied in the presence of increasing amounts of the anion ([SCN-] ) 2 × 10-3-5 × 10-3 mol dm-3]. Solutions with 1:1:25 M:L:SCN- molar ratios were employed to demonstrate the inability of [CoLSCN]+ to bind O2. In such cases the concentration of NaSCN was taken into account in keeping the ionic strength at 0.15 mol dm-3. The equilibrium constants of the anaerobic complexes were determined by means of the computer program HYPERQUAD,23 while a modification of this program was used to calculate the equilibrium constants of the oxygenated complexes from electromotive force data. Such modification was made in order to consider the presence of a chemical species (O2) involved in the complexation equilibria, whose free concentration is kept constant during the measurement. Ligand protonation constants employed in the calculations were determined in a previous work.24 All titration curves were treated either as a single set or as separated entities without significant variation in the value of the stability constants. Results and Discussion Description of the Structure of [NiLCl][NiL(H2O)](ClO4)3. The crystal structure of the compound consists of [NiLCl]+ and [NiL(H2O)]2+ complex cations and perchlorate anions. Figure 1 shows an ORTEP25 drawing of the two complex cations with atom labeling. The coordination environment for each metal is octahedral, determined by three nitrogen and two oxygen atoms of the ligand and completed by secondary species from the medium. In particular, Ni1 is coordinated by N1, N2, N3, O1, and O2 and by a chloride ion (Cl), giving rise to a monocharged cation. The octahedral coordination geometry is rather distorted, as shown by the values of the bond angles reported in Table 2. In fact,
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Table 2. Bond Lengths [Å] and Angles [deg] for [NiLCl][NiL(H2O)](ClO4)3 Ni1-N1 Ni1-N3 Ni1-N2 Ni1-O1 Ni1-O2 Ni1-Cl N1-Ni1-N3 N1-Ni1-N2 N3-Ni1-N2 N1-Ni1-O2 N3-Ni1-O2 N2-Ni1-O2 N1-Ni1-Cl N3-Ni1-Cl N2-Ni1-Cl O2-Ni1-Cl N1-Ni1-O1 N3-Ni1-O1 N2-Ni1-O1 O2-Ni1-O1 Cl-Ni1-O1
2.055(8) 2.094(8) 2.102(7) 2.384(6) 2.156(6) 2.348(3) 100.8(3) 83.3(3) 82.4(3) 78.2(3) 103.4(3) 161.4(3) 101.6(2) 155.8(2) 91.3(2) 89.9(2) 148.2(3) 76.1(3) 126.6(3) 71.9(3) 89.39(18)
Ni2-N4 Ni2-N5 Ni2-N6 Ni2-O3 Ni2-O4 Ni2-O5 O3-Ni2-N4 O3-Ni2-O4 N4-Ni2-O4 O3-Ni2-N6 N4-Ni2-N6 O4-Ni2-N6 O3-Ni2-N5 N4-Ni2-N5 O4-Ni2-N5 N6-Ni2-N5 O3-Ni2-O5 N4-Ni2-O5 O4-Ni2-O5 N6-Ni2-O5 N5-Ni2-O5
2.060(8) 2.098(8) 2.089(8) 2.043(7) 2.067(7) 2.101(9) 157.4(4) 78.5(3) 79.1(3) 82.4(3) 120.3(4) 160.0(3) 101.6(3) 83.7(4) 94.8(3) 83.4(3) 81.6(3) 95.8(4) 92.3(3) 90.6(3) 172.7(4)
the X-Ni-Y bond angular values (with X and Y donors occupying opposite positions in the octahedron) range from 148.2(3)° (N1-Ni1-O1) to 161.4(3)° (N2-Ni1O2), remarkably different from the theoretical expected value (180°). As far as [NiL(H2O)]2+ is concerned, the Ni2 metal ion is coordinated by the N4, N5, N6, O3, and O4 atoms of the ligand and by a water molecule (O5). The coordination geometry is more regular than that of [NiLCl]+, with the X-Ni-Y bond angles falling in the range 157.4(4)-172.7(4)°. In both cations the triamine moiety of the macrocycle is facially coordinated to the metal center. The Ni-N bond distances are very similar (Table 2) in both complexes, while the Ni-O ones are remarkably different, being shorter in [NiL(H2O)]2+. The macrocycle adopts very similar bent conformations in both complex cations. In particular, two almost perpendicular planes are defined by N1, N2, O1, and O2 and N2, N3, and O1 (dihedral angle 101.4(4)°) in [NiLCl]+ and by N4, N6, O3, and O4 and N4, N5, and N6 (dihedral angle 106.4(4)°) in [NiL(H2O)]2+. The most remarkable difference is shown by the exchanged positions occupied by the donor atoms in the two cations. Actually, the previously defined couples of planes intersect each other along the lines N2‚‚‚O1 and N4‚‚‚N6 in [NiLCl]+ and [NiL(H2O)]2+, respectively. In the crystal packing, the two complex cations [NiL(H2O)]2+ and [NiLCl]+ strongly interact each other via H-bond through the coordinated water molecule and the chloride ion (O5‚‚‚Cl, 3.11(1) Å), with the Ni1 and Ni2 metal ions being located 6.553(3) Å apart from each other. Complex Formation under Anaerobic Conditions. As shown by previous works,15,26 the macrocyclic ligand L forms stable complexes with several metal ions, including Ni(II) and Co(II). In the present study, however, we have reinvestigated the formation of such complexes under our experimental conditions. The complex species formed and the relevant stability constants are reported in Table 3. The constants obtained for the [ML]2+ species are in good agreement with previous values,15,26 although under our experimental conditions also mono- and dihydroxo complexes were found. It is noteworthy that,
Table 3. Ligand Protonation and Metal Ion Complexation Constants Determined in 0.15 mol dm-3 NaCl at 298.1 ( 0.1 K reaction
log K
Ni2+ + L ) [NiL]2+ [NiL]2+ + OH- ) [NiLOH]+ [NiLOH]+ + OH- ) [NiL(OH)2]
9.12(2)a 2.4(1)a 3.8(1)a
Co2+ + L ) [CoL]2+ [CoL]2+ + OH- ) [CoLOH]+ [CoLOH]+ + OH- ) [CoL(OH)2]
8.53(2)a 3.2(1)a 4.0(1)a
2[CoL]2+ + O2 + OH- ) [Co2L2(O2)OH]3+ [CoL]2+ a
+
SCN-
)
[CoLSCN]+
11.4(2)b 2.62(8)a
Anaerobic conditions. b Aerobic conditions.
for both [NiL]2+ and [CoL]2+, the equilibrium constants for the addition of the second OH- ion are greater than those corresponding to the binding of the first OH-; namely, [MLOH]+ displays a greater propensity than [ML]2+ to bind hydroxyl ions. This characteristic is rather unusual and may be indicative of the fact that detachment of ligand donor atoms occurs upon coordination of the first hydroxyl ion. Describing the crystal structure of the [NiLCl][NiL(H2O)](ClO4)3 complex, we have evidenced that a significant elongation of the binding contacts between the metal ion and the oxygen donors of the ligand is found in [NiLCl]+ with respect to [NiL(H2O)]2+. Such elongation can be attributed to a weakening of the metal-donors linkage determined by the lower charge density on the metal ion in [NiLCl]+ affecting almost exclusively those bonds (Ni-O) which are mostly electrostatic in nature. As a matter of fact, in the complexes {[ML]3(µ3-CO3)}(ClO4)4 (M ) Cu(II) and Zn(II)), containing the negatively charged carbonate anion, one ligand oxygen atom is not involved in the coordination, while the other one is only weakly coordinated.24 In an aqueous solution, where the dielectric nature of the solvent strongly reduces electrostatic attractions, a larger weakening of the M-O bonds is expected upon OH- binding, supporting the possible detachment of coordinated ligand oxygen atoms. O2 Uptake by [CoL]2+. Recently, we have shown that [CuL]2+ and [ZnL]2+ readily absorb atmospheric CO2, giving rise to {[ML]3(µ3-CO3)}(ClO4)4 species, both in water and in alcoholic solutions, with the reactivity of such complexes toward CO2 being favored by the fact that L is not able to fulfill the coordination sphere of these metal ions.24 On the basis of these results, we were interested in analyzing the reactivity of [CoL]2+ toward molecular dioxygen. Potentiometric titrations performed for the 1:1 Co(II)/L system under an oxygen atmosphere did not display any difference with respect to the anaerobic system in the acidic region, while the alkaline portions of the titration curves accounted for the formation of the dinuclear oxygenated [Co2L2(O2)OH]3+ complex. As shown by the speciation diagram shown in Figure 2, this is the main species in a rather large pH range (8.5-10.5), in solution containing Co(II), L, and O2 in 1:1:1 molar ratios. The assembly of similar dicobalt(II) µ-peroxo-µ-hydroxo complexes from mononuclear entities has already been observed for some other ligands, including openchain13,27-31 and macrocyclic8,11-14 ones. The introduction of OH- bridges gives an important contribution to the stability of similar oxygenated complexes and, in some cases, is crucial for the formation of such species.13
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Acknowledgment This work has been supported by the Ministero dell’Universita` e della Ricerca Scientifica e Tecnologica (MURST, Rome) within the program COFIN 98.
Figure 2. Species distribution diagram for the 1:1:1 system Co(II)/L/O2. Percentages were calculated with respect to Co(II).
Supporting Information Available: Crystal data and structure refinement, atomic coordinates and equivalent isotropic displacement parameters, bond lengths and angles, anisotropic displacement parameters, hydrogen coordinates and isotropic displacements parameters for hydrogen atoms, and least-squares planes. This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited
To further evidence the importance of OH- in promoting the formation of oxygenated complexes, we have inhibited the formation of [CoLOH]+ by competitive coordination with SCN-. As shown by potentiometric titrations, SCN- coordinates to [CoL]2+ (log K ) 2.62; Table 3), and under appropriate conditions (see the Experimental Section), [CoLSCN]+ is the unique species present in solution in a large pH range (7.5-10). Potentiometric measurements performed under an aerobic atmosphere clearly show that this species is insensitive to molecular dioxygen, confirming that the formation of [Co2L2(O2)OH]3+ occurs only in the presence of the µ-OH- bridge. At first glance, this result may seem abnormal. In fact, several studies dealing with the oxygenation of Co(II) complexes of tetraazacycloalkanes (L) have shown that the reactivity of O2 toward such complexes has different features, depending on the presence of cis- or trans-diaquo[CoL(H2O)2]2+ isomers, because only the cis isomer has the appropriate structural requirements for the formation of the dibridged species.7,8,11-14 Hence, the formation of dibridged complexes is favored by smaller tetraazacycloalkanes ([12]aneN4, [13]aneN4), forming cis complexes, in contrast to larger tetraazacycloalkanes ([14]aneN4, [15]aneN4) whose isomeric composition of complexes is almost completely shifted toward the trans species. The formation of dibridged complexes is also inhibited by pentaazacycloalkanes, because only one coordination site is available for the binding of exogen ligands in their Co(II) complexes,9,10 and similarly there are indications that also monooxatetraazacycloalkane ligands9 are not suitable for the formation of µ-peroxoµ-hydroxo species. On the contrary, L forms only the [Co2L2(O2)OH]3+ complex, despite the fact it contains five donor atoms in a large 15-membered macrocyclic ring. The presence in L of oxygen donor atoms, whose coordination properties in water are very sensitive to variation of metal ion charge, and the cofacial coordination of the nitrogen donors, allowing facile achievement of free coordination sites in the cis position, make the [CoL]2+ complex particularly suitable for O2 uptake in the presence of OH- anions. The uncoordinated oxygen atoms of the ligand might assist the formation of the dibridged complexes via intramolecular hydrogen bonds, as proposed for similar species formed by binucleating macrocycles. Unfortunately, we are not able to support this consideration by structural data because all crystals of the complex we have obtained were not suitable for X-ray analysis.
(1) Lindoy, L. F. The Chemistry of Macrocyclic Ligand Complexes; Cambridge University Press: Cambridge, U.K., 1989. (2) Oxygen Complexes and Oxygen Activation by Transition Metals; Martell A. E., Ed.; Plenum: New York, 1988. (3) Basolo, F.; Hoffman, B. M.; Ibers, J. A. Acc. Chem. Res. 1975, 8, 384. McLendon, G.; Martell, A. E. Coord. Chem. Rev. 1976, 19, 1. (4) Dioxygen Activation and Homogeneous Catalytic Oxidation; Simandi, L. I., Ed.; Elsevier: Amsterdam, The Netherlands, 1991. (5) Durand, R. R.; Bencosme, C. S.; Collman, J. P.; Anson, F. C. J. Am. Chem. Soc. 1983, 105, 2710 and references cited therein. (6) Kolchinski, A.; Korybut-Daszkiewicz, G. B.; Rybak-Akimova, E. V.; Busch, D. H.; Alcock, N. W.; Clase, H. J. J. Am. Chem. Soc. 1997, 119, 4160. Collman, J. P.; Herrmann, P. C.; Fu, L.; Eberspacher, T. A.; Eubanks, M.; Boitrel, B.; Hayoz, P.; Zhang, X.; Brauman, J. I.; Day, V. W. J. Am. Chem. Soc. 1997, 119, 3481 and references cited therein. (7) McLendon, G.; Mason, M. Inorg. Chem. 1978, 17, 362. (8) Kodama, M.; Kimura, E. J. Chem. Soc., Dalton Trans. 1980, 327. Machida, R.; Kimura, E.; Kodama, M. Inorg. Chem. 1983, 22, 2055. (9) Kodama, M.; Kimura, E. Inorg. Chem. 1980, 19, 1871. (10) Kimura, E.; Kodama, M.; Machida, R.; Ishizu, K. Inorg. Chem. 1982, 21, 595. (11) Wong, C.-L.; Switzer, J. A.; Balakrishnan, Endicott, J. F. J. Am. Chem. Soc. 1980, 102, 5511. (12) Niedhoffer, E. C.; Timmons, J. H.; Martell, A. E. Chem. Rev. 1984, 84, 137. (13) Cabani, S.; Ceccanti, N.; Tine`, M. R. Pure Appl. Chem. 1991, 63, 1455. (14) Cabani, S.; Ceccanti, N.; Pardini, R.; Tine´, M. R. Polyhedron 1999, 18, 3295. (15) Hancock, R. D.; Bhavam, R.; Wade, P. W.; Boeyens, J. C. A.; Dobson, S. M. Inorg. Chem. 1989, 28, 187. (16) Walker, N.; Stuart, D. D. Acta Crystallogr., Sect. A 1983, 39, 158. (17) Altamore, A.; Cascarano, G.; Giacovazzo, C.; Guagliardi, A. J. Appl. Crytallogr. 1993, 26, 343. (18) Sheldrick, G. M. SHELXL93. Program for the Refinement of Crystal Structures; University of Go¨ttingen: Go¨ttingen, Germany, 1993. (19) International Tables for X-ray Crystallography; Kynoch: Birmingham, England, 1974; Vol. IV. (20) Bianchi, A.; Bologni, L.; Dapporto, P.; Micheloni, M.; Paoletti, P. Inorg. Chem. 1984, 23, 1201. (21) Gran, G. Analyst (London) 1952, 77, 661-671. (22) Oxygen and Ozone, Solubility Data Series; Battino, R., Ed.; Elsevier-Science: New York, 1981. (23) Gans, P.; Sabatini A.; Vacca, A. Talanta 1996, 43, 17391753. (24) Bazzicalupi, C.; Bencini, A.; Bencini, A.; Bianchi, A.; Corana, F.; Fusi, V.; Giorgi, C.; Poali, P.; Paoletti, P.; Valtancoli, C.; Zanchini, C. Inorg. Chem. 1996, 35, 5540.
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(25) Johnson, C. K. ORTEPII; Report-5138; Oak Ridge National Laboratory: Oak Ridge, TN, 1973. (26) Cabral, M. F.; Delgado, R. Helv. Chim. Acta 1994, 77, 515. (27) Cabani, S.; Ceccanti, N.; Conti, G. J. Chem. Soc., Dalton Trans. 1983, 1247. (28) Cabani, S.; Ceccanti, N.; Tine´, M. R. J. Chem. Soc., Dalton Trans. 1988, 373. (29) Cabani, S.; Ceccanti, N.; Conti, G.; Gianni, P. Gazz. Chim. Ital. 1982, 112, 159. (30) Basak, A. K.; Martell, A. E. Inorg. Chem. 1988, 27, 1948.
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Received for review January 24, 2000 Revised manuscript received May 30, 2000 Accepted June 13, 2000 IE000095D
